1896 Inorganic Chemistry, Vol. 12,No. 8, 1973
total of 42 for both 0.10 and 0.30 F ZnC1, is shown in Table IIb, c. These models have been selected so that the probability on the basis of an F test for the model with the highest variance should be higher than 5% as compared with the model with the lowest variance either for the 0.1 F ZnC12 or for the 0.3 F ZnClz solutions. From this table it can be seen that on the basis of the F test it is not possible to distinguish between the best models even if the model involving ZnCL- is the only one with a reasonably high probability at both concentrations. However, if we include the stipulation that for a given model the pK values calculated for each concentration should agree, we find that only one model will fit these criteria, namely, the equilibrium ZnC142 - tj ZnC13- C1- with an average pK value of 2.15 0.04. Furthermore the variance for this system is smaller than the variance found at high pC1- for the KC1-AlC13 system but larger than the variance of 3 X lo-' found at low pC1- for the KCl-AlC1, system (last four measurements in Table Ia) where the reaction mechanism is unambiguous. This is a behavior which one would expect. In view of the unexpected degree of success with which the polymers Zn3Cllo4- and Zn4Cl13'- represent the data, respectively, at 0.10 and 0.30 F ZnC1, it would have been interesting to include n1 values higher than 6 in the above model fitting, but unfortunately this would have required excessive computer time. The ambiguity in choosing a preferred model on the basis of mathematical data fitting is obviously accentuated by the small variation in the degree of zinc coordination provided by the data, as illustrated in Figure 5. If the data had covered a wider range, the probability of the reaction ZnC142- =+-ZnCl; C1- would certainly have been vastly enhanced. However, as matters stand, we are left with trying to predict the total behavior of a function from an imperfect knowledge of only its first few derivatives. Physically there seems to be no way to get around this problem because of the low solubility of ZnC12 at high pC1- values. However, the present measurements
*
+
+
Ronald A. De Marco and Jean'ne M.Shreeve have not been performed all the way up to the saturation point of ZnC1,. It is noteworthy that the model equation 2ZnC1,
'-+ Zn,C1,
3-
t C1-
proved thoroughly inadequate as a representation of the data. This equation has been p r o p o ~ e d ~ as ~ ?the ~ ' best representation of the structural situation in MCl-ZnCl, mixtures rich in MC1 but it is,a very poor representation of the buffering action of ZnCl, in KC1-AlC13 mixtures. If we suppose that the structural reality behind the buffering action of ZnC1, is correctly represented by Z n C 4 2- =+ ZnC13- C1-, we can reconcile this behavior with that in MCl-ZnCl, through the additional supposition that ZnC13is four-coordinate with the fourth chloride forming a bridge to aluminum so as to give an entity wch as C13ZnC1AlC13'-. Such an entity would, of course, be impossible in MCl-ZnC12 melts while perhaps being favored in the present case because of the high concentration of chloroaluminate species. The fact that this concentration is so high and, therefore, so slightly affected by solvation reactions obviously precludes the possibility of detecting solvation reactions by means of activity measurements.
+
Registry No. KCl, 7447-40-7;AlCl3, 7446-70-0;ZnClz, 7646-85-7; A1cl4-, 17611-22-2;ZnC14'-, 15201-05-5. Acknowledgment. The authors wish to thank B. LottrupKnudsen for help with programming and F. W. Poulsen and G. P. Smith for valuable discussions. Further thanks are due to Statens tekniskvidenskabelige Fond for financial support for J. H. von Barner and to the Northern Europe University Computing Center (NEUCC) for computing time. (24) C. A. Angel1 and D. M. Gruen, J . Phys. Chem., 70, 1601 (1966). (25) W. E. Smith, J. Brynestad, and G. P. Smith, J . Chem. Phys., 5 2 , 3 8 9 0 (1970).
Contribution from the Department of Chemistry, University of Idaho, Moscow; Idaho 83843
Reactions of the Difluoramine-Potassium Fluoride Adduct, HNFz KF, with Sulfinyl and Perfluoroalkylsulfinyl Fluorides. Preparation of Perfluoroalkyl Perfluoroalkylthiosulfonates RONALD A. De MARC0 and JEAN'NE M. SHREEVE*
Received January 22, 1 9 73 The reaction of HNF,.KF with sulfinyl fluoride gave FS(O)NF, which was isolated and had sufficient stability t o be characterized by spectroscopic methods. Attempts to prepare CF,S(O)NF, and C,F,S(O)NF, resulted in the formation Efforts t o synthesize of N,F, and the perfluoroalkyl perfluoroalkylthiosulfonate esters CF,SO,SCF, and C,F,SO,C,F,. additional thiosulfonate esters led to the characterization of C,F,S(O)Cl, "F nmr spectral identification of C,F,SO,SCF, and CF,SO,SC,F,,and spectroscopic characterization of the previously reported C,F,SO,F.
Introduction In surveying the field of fluorinated sulfur-nitrogen compounds, the dearth of compounds which contain the difluoroamino moiety bound t o sulfur became apparent. With the exception of CF3SNF2,' the known difluoroamino (1) Alfred P. Sloan Foundation Fellow, 1970-1972. (2) E. C. Stump and C. D. Padgett, Iflovg. Chem., 3, 6 1 0 ( 1 9 6 4 ) .
derivatives of this class of compounds are limited t o cases in which sulfur is present in its maximum oxidation state.3 Photolytic or thermal reactions of NzF4 for the preparation of sulfur(1V) difluoroamino compounds were not useful since oxidation of sulfur occurred with SF4 derivatives or (3) S. M. Williamson, Progr. Inorg. Chem., 7, 39 ( 1 9 6 6 ) .
Inorganic Chemistry, "01. 12, No. 8, 1973 1897
HNFz~KFReactions with SO2 , 3 while with sulfinyl halide^,^ no sulfur-nitrogen derivatives were isolated, e.g. RfSF,
+ N,F,
SO,*+ N,F, SOCl,
.+.
-+
RfSF,NF,
Rf = CF,, F
FSO,NF,
+ N2F4-+ NF,C1
SOBr, t N,F,
.+.
N,F,
Generally the reactions of HNFz involved the use of strong acids or Lewis bases. For example, recently it was demonstrated that the NF2 group can be readily introduced through the use of an HNFz-KF adduct. Thus, perfluoroamides or -amines were prepared in high yield with this adduct under mild reaction conditions' s6 -23'
RfC(0)F + HNF,.KF -RfC(O)NF, CF,(OF),
-80'
+ HNF,.KF -CF,(OF)ONF,
+ CF,(ONF,),
The ease with which the HNFzaKF adduct reacted suggested its potential application in the preparation of sulfur(1V)difluoroamino derivatives. Results and Discussion Difluoroaminosulfinyl fluoride, FzNS(O)F, was obtained in low yield from the reaction of SOFz with the HNFZ-KF adduct. The compound was unstable in the gas phase and in solution and, therefore, was usually isolated with varying amounts of sulfur dioxide as the major impurity. Although only minor amounts of FzNS(0)F were isolated, the reaction is believed to proceed in better than 75% yield based on the initial HNFz present. The major nitrogen-fluorine-containing compound isolated was N2F4 which almost certainly must arise from the decomposition of the sulfinamide and not from decomposition of HNFz. Although it was reported that HNF2 was converted to NzF4 in high yield by such catalysts as stainless steel, copper, urea, and LiH after a contact time of more than 68 hr at temperatures greater than 0°,7aa subsequent paper showed that HNFz in contact with KF at 25" decomposed to N2F2 exclusively.% When HNFz was allowed to contact urea at -go", the pressure decreased but on warming to 25" all the HNFz was recovered. In our previous work with HNF2*KFperfluoroacyl fluoride reactions at -23' and below, no trace of N2F4 was observed.' Since in our present study, reactions were carried out in glass vessels at -78" with reaction times of 6-12 hr, there is no reason to expect the HNFz to behave in a different manner. Although it is not possible to exclude the possibility completely that the sulfur compounds could act as decomposition catalysts, the NtF4 is a likely product in the decomposition of FzNS(0)F because of the weak sulfur-nitrogen bond. The fate of the FS(0). radical, which would be expected to form, is not clear. The decomposition of this radical is much more complex as is suggested by the formation of solids ranging from yellow to violet and red during the reaction and nmr spectral runs. A sample of FzNS(0)F which was allowed to decompose in an infrared cell resulted in the formation of SOZ and S0F2. Both of these compounds would be expected in FS(0). decomposition and disproportionation. The N-F stretching (4) L. M. Zaborowski, K. E. Pullen, and J. M. Shreeve, Inorg. Chem., 8 , 2005 ( 1 9 6 9 ) . (5) R. A. De Marco and J . M. Shreeve, Inorg. Chem., 10, 911 (1971). (6) D. Pilipovich and M. G. Warner, U. S. Patent 3 , 6 6 3 , 5 8 8 (May 16, 1972); Chem. Abstr., 7 7 , 100747 (1972). ( 7 ) (a) E. A. Lawton and J. Q. Weber, J. Amer. Chem. SOC.,8 5 , 3595 ( 1 9 6 3 ) ; (b) E. A. Lawton, D. Pilipovich, and R. D. Wilson, Inorg. Chem., 4, 1 1 8 (1965).
region in the infrared spectrum was too complex to readily assign N2F4 as a product but "F nmr spectral studies do confirm the formation of SOFz and N2F4 as the decomposition products in addition to some N2F2 which can also arise from NFz radicals. The stability of FzNS(0)F varies with its physical state. Essentially complete decomposition in the gas phase occurred in about 8 hr, while complete decomposition in an SOZFZ solution occurred within 1 hr. The very low recovery and poor reproducibility of this reaction prevented complete characterization of F*NS(O)F, but the spectral properties readily confirmed the assigned structure. The infrared spectrum consists of bands at 1305 (s, d), 925 (sh, br), 905 (s), 885 (sh, br), 860 (s), and 765 (m) cm-'. The complex region between 925 and 860 cm-' can be assigned to the asymmetric and symmetric NF2 stretching frequencies. The doublet at 1305 cm-' corresponds to the S=O stretch and the band at 765 cm-' to the S-F stretch. In comparison the S=O stretching frequencies for SOFz and SOFCl are doublets centered at 1330 and 1290 cm-', respectively. The S-F stretch in SOFCl is at 750 cm-'. The "F nmr spectrum of F2NS(0)F consists of a triplet at q3 -19.6 (JNF,-sF = 9.5 Hz) and a typically broadened NF2 resonance at q3 -29.8 with the area ratio 1 : 1.9. The reactions of perfluoroalkylsulfinyl fluorides with HNFz.KF were carried out in an attempt to prepare stable N-fluorosulfinamide derivatives. The reactions are comparable to the sulfinyl fluoride reaction in that N2F4 was the major product isolated but no corresponding sulfinamide was found. The only sulfur-containing compound isolated from both CF3S(0)F and CzFSS(0)Fwas the symmetrical perfluoroalkyl perfluoroalkylthiosulfonate, RfS02SRf. These esters could arise from the interaction of the corresponding RfS(0). radicals 2RfS(O)NF,
-+
2RfS(O). + N,F,
4 RfS0,SRf
Isolation of these esters as very stable compounds suggests that the difference in products observed for the decomposition of NF2S(O)F and the hypothetical RfS(O)NF2 is due to the greater stability of the RfS(O)*radical compared to that of FS(0). which permits dimerization and rearrangement to occur. The unfluorinated sulfinamides (RS(O)NR2) and trifluoromethyl sulfinamides (CF3S(0)NR2) are stable to 1OO', but with the introduction of the electron-withdrawing NFz group, in addition to the perfluoroalkyl group or fluorine, the net effect is to weaken the sulfur-nitrogen bond. For NF2S(O)F, the yield of NZF4 accounted for about 80% of the HNF2 initially consumed in the reaction. The yields of the thiosulfonate esters were more variant and although the trifluoromethyl derivative accounted for 80% of the NzF4, the pentafluoroethyl compound gave only 30% of the ester based on N2F4. In order to obtain sufficient amounts of CzF5S02SC2FS for complete characterization, other less hazardous methods were attempted. The most direct method to prepare compounds of the RfSOzSRf type is by the reaction of the corresponding sulfinyl chloride with mercury.8 Although this reaction gave good yields for CF3S02SCF3, it is not very effective for the formation of C2FSSO2SCzF5. The two main products were CzFSSOzCland (C2FSS)2. The pentafluoroethyl pentafluoroethylthiosulfonate was isolated (8) D. T.
Sauer, University
of Idaho, private
communication.
Ronald A. De Marco and Jean'ne M. Shreeve
1898 Inorganic Chemistry, VoZ. 12, No. 8, 1973
in quantities to enable spectroscopic identification only. The infrared spectrum consisted of bands at 1425 (ms), 1330 (sh, m), 1320 (ms), 1242 (s), 1140 (ms), 1128 (sh, m), 977 (sh, m), 952 (ms), 756 (m), 535 (ms), 545 (w), and 524 (w) cm-' . The 19Fnmr spectrum contained four resonances which are listed in Table I. The triflu~romethyl~ and pentafluoroethyl thiosulfonate esters are the only reported perfluoro examples of this class of compounds which contain sulfur in different oxidation states. Since each contains only one kind of perfluoroalkyl group, it was of interest t o attempt the preparation of unsymmetrical analogs for spectroscopic comparison. Unfortunately, neither the reaction of a mixture of CF3S(0)C1 and C2F5S(0)Cl with mercury nor the reaction of (CF3S(0)O)2Zn with C2F5SC1was consistent or occurred in sufficiently high yield to allow such a comparison. The reaction of equimolar amounts of the sulfinyl chlorides with mercury did result in the formation of a mixture of two compounds which, after purification via gas chromatography: were studied with 19F nmr. Based on nmr spectral shifts and coupling constants compared with the symmetrical CF3S02SCF3 and C2F5S02SC2FS and other CF3 and C2F5 sulfonyl and sulfenyl derivatives, the mixture was shown to contain CF3S02SCZFS and C2F5S02SCF3(Table I). As can be seen these values are in good agreement with reported values. The "F nmr spectra of sulfinyl derivatives which contain the -CF*S(O)- group have been interpreted as having nonequivalent methylene fluorines due to the asymmetric center at the sulfur Although the methylene fluorines in both CF3CF2S(0)F and CF3CF2S(O)Cl exhibit this non. equivalence, the effect is more pronounced in the latter case. Since the presence of an asymmetric center should not account for this difference, it is possible that a rotational factor also contributes. Experimental Section Apparatus. Volatile liquids and gaseous materials were handled in a standard Pyrex vacuum line equipped with a Heise Bourdon tube gauge. The reactions were carried out in 25-1111 Pyrex flasks equipped with Fischer-Porter stopcocks. The products were purified by fractional condensation or gas chromatography with 0.25-in. copper columns packed with FS-1265 or Kel-F No. 3 oil on Chromosorb P. Analysis. Infrared spectra were taken with a Perkin-Elmer Model 457 infrared spectrometer using a 10-cm Pyrex glass cell equipped with KBr windows or a 10-cm stainless steel cell equipped with AgCl windows. High-resolution I9F nmr spectra were obtained with a Varian Model HA-100 spectrometer operating at 94.1 MHz. Fluoride analyses were done by sodium fusion and the fluoride ion determined by the use of an Orion specific ion electrode Model 9409A or by Beller Mikroanalytisches Laboratorium, Gottingen, Germany. Starting Materials. Difluoramine was prepared from trityldif l ~ o r a m i n e '(PCR, ~ Inc.) and was complexed as previously described.' Sulfinyl fluoride was prepared by exposing SF, (K & K Laboratories) to moist air in a Pyrex vessel and then isolated by fractional condensation. Sulfuryl fluoride was prepared by the reaction of SO, and fluorine. Trifluoromethyl- and pentafluoroethylsulfinyl fluoride^,'^ trifluoromethylsulfuryl chloride, and zinc trifluoromethylsulfonate' were prepared by the literature methods. Diglyme was purchased from Aldrich Chemical Co. and stored over (9) R. N. Haszeldine and I . M. ICidd, J. Chem. Soc., 2 9 0 1 (1955). (10) D. T. Sauer and J . M. Shreeve, J. Fluorine Chem., 1, (1 971). (11) J. I. Darragh, A. M. Noble, D. W. A. Sharp, D. W. Walker, and J . M. Winfield,Inorg. Nucl. Chem. L e f t . , 4, 517 (1968). ( 1 2 ) F. Sed, J . Boudier, and W. Gombler, Chem. Ber., 102,4 4 3 (1969). (13) W. H. Graham and C. 0. Parker, J. Org. Chem., 2 8 , 8 5 0 (1963). (14) C. T. Ratciiffe and J . M. Shreeve, J. Amer. Chem. Soc., 90, 5403 (1968).
Table I. I9F Nmr Spectraa
~
5
-
CF ,-SO Z-SCF, 76.8
36.3
1
r 7 . 6 1 ~ 2 . 9 1 i-1.11r-7.2-1 CF ,-SO,-SCF,-CF CF3-CF,-SO ,-SCF, L 0 . 7 - A 0 . 4 d 77.2
87.7
83.4
77.9
77.8 111.1 87.3 83.4 80.5 113.5 -44.8
111.5
78.4
36.4
116.4
J C F ~ - F = 0.95, JCF3-F' =
1 . 0 5 , J ~ -= ~f 23 1
125.0 F /
CF,C-S(0)F \
F' 79.6 124.7 19.4 JCF,-F = 1.1, JCF,-F' = 1.0, J F - F ~= 5.2, J s F - F ' = ~ . ~J ,C F ~ - S F = 7.7 0
Chemical shifts in @ (internal CC1,F reference); Jvalues in hertz.
molecular sieves. Potassium fluoride (J. T. Baker) was dried at 350" and CsF (American Potash and Chemical Corp.) was dried in vacuo at 150". Tetrafluoroethylene and pentafluoroethyl iodide (Hynes Chemical Co.) were used as received and the bis(pentafluor0ethyl) disulfide was prepared via the published method.15 Warning! Care should be exercised in handling HNF, since it is known t o exhibit explosive properties especially when condensed at -183 or -196". Liquid HNF, is more shock sensitive than nitroglycerine and should be handled in small quantities. Adequate shielding should be used at all times including the cleaning of reaction vessels since in one case the exposure of t h e solid residue t o air resulted in a violent explosion. Reaction of SOF,. Approximately 1 8 mmol of SOF, was condensed at -138" onto 4.0 mmol of HNF,.KF and then warmed to -78" for 6-12 hr. The volatile products were removed at -78", then at -23", and finally at 25". The difluoroaminosulfinyl fluoride was isolated usually by fractional condensation in a trap cooled t o -108" and was invariably contaminated with SO, and occasionally with some HNF, (0.1-0.2 mmol). All attempts t o purify the sample further by physical or chemical methods resulted in further decomposition. Also isolated from the reaction mixture were unreacted SOF, and HNF, which were then recylced. It was impossible to determine the amounts of the two compounds since they were both trapped at -128" but HNF, was present in less than 1-mmol amounts. Approximately 1.6 mmol of N,F, was also found. Reaction of CF,S(Q)F. This reaction was carried out under essentially the same conditions as with SOF, (18 mmol of CF,S(OIF-4 mmol of HNF,.KF). Fractional condensation with traps at -91, -138, and -196" resulted in the isolation o f 1.4 mmol of CF,SO,SCF, (identified by the reported infrared' and "F nmr spectra), unreacted starting materials and 1.8 mmol of N,F,. Preparation of C,F,S(Q)CL. A mixture of 13.8 mmol of C,F,S(O)F and 16.0 mmol of HC1 was warmed slowly t o 25" in a 1-1. Pyrex vessel and allowed to remain at this temperature for 45 min. After fractional condensation at -78", the sample was purified by gas chromatography using a 17-ft column packed with 20% FS-1265 on Chromosorb P. The infrared spectrum consists of bandsat 1340 (m), 1250 (s), 1228 (s), 1140 (m), 943 (m), 753 (mw), 505 (mw), and 478 (m) cm-'. A molecular weight of 203.7 (202.6) was found and the fluoride analysis yielded 45.7% (46.1% required). From vapor presswe data represented by the equation logPTorr = 7.974 - 1707.4/T°K (valid in the range 0-65"), a boiling point of 62.2' was obtained. The heat of vaporization was calculated to be 7.8 kcal/mol and the Trouton constant is 23.3 eu. "F nmr data for both C,F,S(O)F and C,F,S(O)Cl are listed in Table I . Preparation of C,F,SQ,F. A mixture of 20 mmol of C,F4 and 46 mmol of SO,F, was condensed into a 30-ml Hoke bomb which (15) W. J. Middleton, U. S. Patent 3,069,395 (Dec 18, 1962); Chem. Abstr., 59, P1493h (1962).
Inorganic Chemistry, Vol. 12,No. 8, 1973 1899
Platinum(I1) Chloride-Aluminum Chloride Complex contained dried CsF and about 7 ml of diglyme. After heating the bomb at 80' for 4 days, the volatile portion was fractionated (-98') and then further purified by gas chromatography by using an 8-ft column packed with 20% Kel-F oil on Chromosorb P. The pentafluoroethylsulfuryl fluorideI6 (30%) was further characterized by 19Fnmr spectra (Table I) and by its infrared spectrum which has bands at 1476 (vs), 1336 (m), 1255 (vs), 1240 (sh, vs), 1208 (m), 1152 (s), 999 (m), 817 (s), 754 (w), 644 (sh, w), and 610 (sh, w) cm-'. Its experimental molecular weight was 203 (202). Preparation of C,F,SCl. A mixture of (C,F,S), (4.75 mmol) and C1, (6.0 mmol) was heated at 90-100" for 6 hr in a Pyrex vessel. Separation of the mixture by fractional condensation gave unreacted disulfide (-78"), C,F5SCl(-108"), and a mixture of C1, and C,F,SCl at -196'. The infrared spectrum consists of bands at 1330 (m), 1230 (vs), 1168 (m), 1122 (s), 980 (s), 965 (sh), 753 (m), and 550 (w)cm-'. Reaction of C,F,SCl and (CF,S(O)O),Zn. Approximately 1 mmol of C,F,SCl was condensed onto (CF,S(O)O),Zn in a Pyrex vessel and the volatile products were monitored every 6 hr over a 36-hr period. The major products identified were CF,C1, (C,F,S),, C,F,SCF,,'" and some anreacted C,F,SCl. No CF,SO,SC,F, was detected during this period. Preparation of CF,SO,SC,F, and C,F,SO,SCF,. A mixture of 5 mmol each of CF,S(O)Cl and C,F,S(O)Cl was condensed at -183" (16) T. (195'9).
Gramstad and R. N. Haszeldine, J. Chem. SOC.,2640
into a 500-ml Pyrex flask equipped with a Teflon stopcock and then warmed to 25'. Over a period of 1.5 hr approximately 3.5 ml of mercury was added with shaking. The mixture was fractionated through traps cooled t o -63, -78, and -196'. Further purification of the materialsretained in the traps at -63 and -78" was accomplished with gas chromatography using an 8-ft column packed with Kel-F No. 3 oil on Chromosorb P. This resulted in the isolation of 0.4 mmol of a mixture of the thiosulfonates. Additional attempts t o separate the mixture by gas chromatography were unsuccessful and attempts t o reproduce or alter this reaction resulted in no isolation of the two products.
Registry No. HNF2 .KF, 13841-05-9;SOF?, 7783-42-8; FS(O)NFz, 39937-04-7; CF3S(O)F, 812-12-4;CF3SOZSCF3, 358-15-6; CZFSS(O)F, 20621-3 1-2; HCl, 7647-01-0; CZFSS(O)C1,39937-08-1; CZF4, 116-14-3; SOZF2,2699-79-8; CzFSSOzF, 354-87-0;(CZFSS)2, 679-77-6; Clz, 7782-50-5; C2F,SC1, 39937-11-6; [CF3S(0)0I2Zn, 39971-65-8; CF3S(O)C1,20621-29-8; CF3S02SCZF5,39937-13-8;C2F5SOzSCF3, 39937-14-9; Hg, 7439-97-6; C2FSSO2SCTF5, 3993715-0. Acknowledgment. Fluorine research at the University of Idaho is supported by the National Science Foundation and the Office of Naval Research.
Contribution from Ames Laboratory-USAEC and the Department of Chemistry, Iowa State University, Ames, Iowa 50010
Spectrophotometric Study of the Platinum(I1) Chloride-Aluminum Chloride Vapor Complex G . N. PAPATHEODOROU' Received February I , 1973
The reaction of solid cluster platinum(I1) chloride (Pt6C1,,) with gaseous aluminum chloride to form a purple gaseous complex has been studied spectrophotometrically. Thermodynamic considerations suggest the stoichiometry for the reaction: '/6Pt6Ci,,(s) A1,Cl6(g) + PtAl,Cl,(g) (hHR = 7.8 kcal/mol, ASR = 6.4 eu). The visible electronic absorption spectrum of the gaseous complex was compared with the spectrum of Pt(I1) centers in liquid LiCl and interpreted in terms of a square-planar PtCl, group sharing edges with two tetrachloroaluminates.
+
Introduction that gaseous aluminum chloride reacts It has been with many MClk metal chlorides to form gaseous complexes with the general formula (MClh)m(AlZC16)n.In most cases studied m = 1 while the value of n varies between and 2. Among the different metal chlorides investigated by vapor pressure equilibrium measurements were some alkali and alkaline earth chlorides (M = Na, K, Mg, Ca)233 and some chlorides of 3d metal cations (M = Mn(II), Fe(II), Co(II), Ni(II), Zn(II)).3 The chlorides of Nd(III),4a U(III), U(V)?b and Pd(II)' have been studied spectrophotometrically. In the present report the existence of a purple vapor complex of platinum chloride with aluminum chloride has been (1) Ames Laboratory USAEC postdoctoral fellow, 1970-1971. Present address: The James Franck Institute, University of Chicago, Chicago, Ill. 60637. (2) (a) E. W. Dewing, Nature (London), 214, 483 (1967); ( b ) M e t .
Trans., 1, 2169 (1970). (3) K. N. Semenenko, T. N. Naumova, L. N. Gorokhov, and A. V. Bovoseloca, Dokl. Akad. Nauk SSSR, 154, 649 (1964). (4) (a) H. A. Oye and D. M. Gruen, J. Amer. Chem. SOC.,91, 2229 (1969); (b) D. M. Gruen and R. L. McBeth, Inorg. Chem., 8, 2625 (1969). ( 5 ) G. N. Papatheodorou, J. Phys. Chem., 77, 472 (1973). (6) K. Lascelles and H. Schafer, 2. Anorg. ANg. Chem., 383, 249 (1971).
established. Furthermore, the equilibrium (m/6)Pt6 c1,z (S) 4- nA12C16 (g)
--f
(PtC1, )m(A&Cls)n(g)
(1)
has been studied by measuring spectrophotometrically the partial pressures of the gaseous complex(es) over the "cluster" platinum chloride (Pt6ClI2). The thermodynamic quantities of the above equilibrium and the electronic absorption spectrum of the gaseous complex are discussed in terms of the stoichiometry and possible structure of the gaseous molecule. Experimental Section Chemicals and Equipment. High-purity anhydrous aluminum chloride was prepared from Baker Analytical reagent AlCl, by slow sublimation in silica tubes under vacuum. The sublimation was repeated six times. The platinum(I1) chloricte was prepared as previously described.'* Large quantities of the cluster platinum chloride Pt6C1,, were prepared by recrystallization of PtC1, from its dilute solutions in liquid Al,Cl, .7b The spectrophotometric measurements were performed on a Cary Model 14 spectrophotometer equipped with a high-temperature cell compartment* (15-cm maximum path length). The quartz cells were the uv-type cylindrical cells, purchased from Pyrocell. The (7) (a) G. N. Papatheodorou and G. P. Smith, J. Znorg. Nucl. Chem., 3 5 , 799 (1973); (b) G. N. Papatheodorou, unpublished work. (8) R. A. Lynde and J. D. Corbett, Inorg. Chem., 10, 1746 (1971).
